Acids and Bases: Properties, Definitions, and Reactions

Acids and Bases

Introduction

This chapter introduces the fundamental concepts of acids and bases, their properties, common examples, and the chemical reactions in which they participate. Understanding acids and bases is essential for studying chemical reactions, solution chemistry, and many biological and industrial processes.

Properties of Acids

General Properties

• Sour taste: Many acids taste sour (e.g., citric acid in lemons).

• Dissolve many metals: Acids can react with and dissolve certain metals, producing hydrogen gas.

• Turn blue litmus paper red: Acids cause a color change in litmus paper due to their effect on pH-sensitive dyes.

Examples and Applications

• Hydrochloric acid (HCl): Found in stomach acid, used in industry for cleaning metals and processing foods.

• Sulfuric acid (H2SO4): Most commonly produced chemical in the US, used in fertilizers, batteries, and manufacturing.

• Nitric acid (HNO3): Used in fertilizers, explosives, and dyes.

• Acetic acid (CH3COOH): Gives vinegar its sour taste; an example of a carboxylic acid (contains the COOH group).

• Carboxylic acids: Found in many living organisms (e.g., citric acid in citrus fruits, malic acid in apples).

Properties of Bases

General Properties

• Bitter taste: Bases often taste bitter (e.g., caffeine in coffee).

• Slippery feel: Bases react with oils on the skin to form soap-like substances, giving a slippery sensation.

• Turn red litmus paper blue: Bases cause a color change in litmus paper due to their effect on pH-sensitive dyes.

Examples and Applications

• Sodium hydroxide (NaOH) and potassium hydroxide (KOH): Used in laboratories, soap, and plastic manufacturing; active ingredient in drain cleaners.

• Sodium bicarbonate (NaHCO3): Baking soda; used as an antacid to neutralize stomach acid.

• Alkaloids: Organic bases found in plants, often bitter and sometimes poisonous (e.g., coniine in hemlock).

Definitions of Acids and Bases

Arrhenius Definition

• Acid: Produces H+ ions in aqueous solution.

• Base: Produces OH- ions in aqueous solution.

Example: HCl is an Arrhenius acid because it ionizes in water to form H+ and Cl- ions.

$\mathrm{HCl(aq) \rightarrow H^+(aq) + Cl^-(aq)}$

In water, H+ ions associate with H2O to form hydronium ions:

$\mathrm{H^+ + H_2O \rightarrow H_3O^+}$

NaOH is an Arrhenius base because it dissociates in water to form Na+ and OH- ions:

$\mathrm{NaOH(aq) \rightarrow Na^+(aq) + OH^-(aq)}$

Brønsted–Lowry Definition

• Acid: Proton donor (donates H+).

• Base: Proton acceptor (accepts H+).

Example: HCl donates a proton to water:

$\mathrm{HCl(aq) + H_2O(l) \rightarrow H_3O^+(aq) + Cl^-(aq)}$

NH3 accepts a proton from water:

$\mathrm{NH_3(aq) + H_2O(l) \rightleftharpoons NH_4^+(aq) + OH^-(aq)}$

Amphoteric Substances

• Some substances, such as water, can act as either acids or bases. These are called amphoteric substances.

Conjugate Acid-Base Pairs

In Brønsted–Lowry acid-base reactions, acids and bases form conjugate pairs:

• Conjugate acid: Formed when a base gains a proton.

• Conjugate base: Formed when an acid loses a proton.

Example: NH3 and NH4+ are a conjugate acid-base pair.

Reactions of Acids and Bases

Neutralization Reactions

When an acid reacts with a base, they neutralize each other to form water and a salt.

$\mathrm{H^+(aq) + OH^-(aq) \rightarrow H_2O(l)}$

General equation: Acid + Base → Water + Salt

Example: HCl(aq) + KOH(aq) → H2O(l) + KCl(aq)

Acid Reactions with Metals and Metal Oxides

• Acid + Metal: Produces hydrogen gas and a salt.

• Acid + Metal Oxide: Produces water and a salt.

Example: 2HCl(aq) + Mg(s) → H2(g) + MgCl2(aq)

2HCl(aq) + K2O(s) → H2O(l) + 2KCl(aq)

Acid-Base Titration

Principle

• Titration is a laboratory technique used to determine the concentration of an unknown acid or base by reacting it with a solution of known concentration.

• The equivalence point is reached when the amount of acid equals the amount of base, as determined by the reaction stoichiometry.

Example: Titrating HCl with NaOH:

$\mathrm{HCl(aq) + NaOH(aq) \rightarrow H_2O(l) + NaCl(aq)}$

Strength of Acids and Bases

Strong Acids and Bases

• Strong acids: Completely ionize in solution (e.g., HCl, HNO3, H2SO4).

• Strong bases: Completely dissociate in solution (e.g., NaOH, KOH).

Weak Acids and Bases

• Weak acids: Only partially ionize in solution (e.g., acetic acid, HF).

• Weak bases: Only partially accept protons or produce OH- (e.g., NH3).

Self-Ionization of Water and the pH Scale

Self-Ionization of Water

Water can act as both an acid and a base, undergoing self-ionization:

$\mathrm{2H_2O(l) \rightleftharpoons H_3O^+(aq) + OH^-(aq)}$

At 25°C, the ion product constant for water is:

$K_w = [H_3O^+][OH^-] = 1.0 \times 10^{-14}$

pH Scale

• pH is a measure of the hydrogen ion concentration in a solution.

• pH < 7: Acidic solution

• pH = 7: Neutral solution

• pH > 7: Basic solution

The pH is calculated as:

$\mathrm{pH = -\log[H_3O^+]}$

To find [H3O+] from pH:

$\mathrm{[H_3O^+] = 10^{-pH}}$

Buffers

Definition and Function

• Buffer: A solution that resists changes in pH when small amounts of acid or base are added.

• Contains significant amounts of both a weak acid and its conjugate base.

• Important in biological systems (e.g., blood maintains pH between 7.36 and 7.40).

Example: Acetic acid and sodium acetate buffer system:

• Added base: Acetic acid neutralizes it.

• Added acid: Acetate ion neutralizes it.

Summary Table: Properties of Acids and Bases

Property	Acids	Bases
Taste	Sour	Bitter
Feel	No characteristic feel	Slippery
Litmus Test	Turns blue litmus red	Turns red litmus blue
Reaction with Metals	Often dissolves metals, produces H2 gas	Usually no reaction
Examples	HCl, H2SO4, CH3COOH	NaOH, KOH, NH3